The_Story_of_an_Iron_Nail curious facts that everyone should know.pptx

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1. ഇരുമ്പാണിയുടെ കഥ പൗരാണിക കാലം മുത ൽ ആധുനിക രസതന്ത്രം വരെ

2. Introduction An iron nail might seem simple, but it carries a long history of technological evolution and cultural meaning.

3. Ancient Origins Nails were used as early as 3400 BCE in Ancient Egypt. Romans perfected iron nails for construction.

4. Medieval Uses By the Middle Ages, blacksmiths commonly forged nails by hand for use in architecture and shipbuilding.

5. The Industrial Revolution Machine-cut nails began replacing hand-forged ones in the 18th century, revolutionizing construction.

6. Wire Nails and Mass Production By the 19th century, wire-drawn nails became dominant due to ease and speed of manufacturing.

7. Types of Nails Box nails, finishing nails, masonry nails—each type tailored to different structural and material needs.

8. Early Nail making Traditional methods involved heating iron rods and shaping them with hammer and anvil. Blacksmiths heat iron ore with carbon to form a dense mass of metal, which is then placed into the shape of square rods and left to cool.

9. Carbon is used to reduce iron ore (usually iron oxides like hematite Fe₂O ₃ or magnetite Fe₃O ₄) into metallic iron. Iron ore itself cannot be used directly in forging or shaping because it’s in a chemically bound, oxidized form. Carbon acts as a reducing agent , removing the oxygen from iron oxide. Fe₂O ₃ + 3C → 2Fe + 3CO ↑ Besides reduction, small amounts of carbon can enter the iron structure, forming steel . Carbon strengthens iron and changes its properties significantly: Wrought iron : very low carbon content Steel : moderate carbon (0.2–2.0%) Cast iron : high carbon (~4%) Why Carbon

10. Fe₂O ₃ + 3C → 2Fe + 3CO ↑ Enters Thermodynamics This is a redox reaction where carbon reduces the iron oxide. The thermodynamic feasibility of this reaction is determined by the Gibbs free energy change (ΔG) : If ΔG < 0 , the reaction is spontaneous. At high temperatures (~900–1200°C), ΔG becomes negative, making the reduction thermodynamically favorable. Carbon becomes a better reducing agent at higher temperatures. CO is more stable than CO₂ at high T, facilitating oxygen removal from iron oxide.

11. Iron Chemistry Basics Iron reacts with oxygen and water to form rust: 4Fe+3O 2 ​+6H 2 ​O → 4Fe(OH) 3 ​→ Fe 2 ​O 3 ​⋅xH 2 ​O (rust) . This is a redox reaction involving: Oxidation of iron (Fe → Fe²⁺/Fe³⁺) Reduction of oxygen (O₂ → OH⁻ in water)

12. How is r ust formed ? H 2 O H 2 O H 2 O H 2 O H 2 O H 2 O Fe 2+ Fe 2+ Fe 2+ Fe 2+ Fe 2+ Fe 2+ H 2 O(l) + e - ⇌ ½H 2 (g) + OH - (aq) -0.83 V Fe 2+ (aq) + 2e - ⇌ Fe(s) -0.45 V Small circle = holds on to electrons more strongly Larger circle = holds onto electrons more weakly Excess electrons – negative charge Positively charged ion in solution ATTRACTION

13. Fe 2+ Fe 2+ Fe 2+ OH - OH - H 2 O H 2 O H 2 O OH - OH - ½ O 2 H 2 O H 2 O H 2 O Fe 2+ Fe 2+ Fe 2+ H 2 O H 2 O H 2 O OH - OH - ½ O 2 H 2 O OH - OH - ½ O 2 H 2 O Fe 2+ (aq) + 2e - ⇌ Fe(s) -0.45 V ½O 2 (g) + H 2 O(l) +2e - ⇌ 2OH - (aq) +0.40 V ½ O 2 What needs to happen next … ? Oxygenated water near surface – able to take electrons from iron Oxygenated water accepts electrons Hydroxide ions formed Iron ions dissolve How is rust formed?

14. Electrochemical Process of Rust All electrochemical processes involve a type of chemical reaction called oxidation-reduction or redox reactions. In a redox reaction, there is a transfer of electrons. The transfer of electrons in corrosion are taken from the surface of the metal and transferred to suitable electron acceptors, like oxygen and hydrogen. The anode will be a location where the metal is stressed or damaged. Cathode is another part of the metal not undergoing corrosion. Water acts as the electrolyte. At the anode, oxidation half-reaction: Fe(s) → Fe 2+ ( aq ) + 2e- At the cathode, a reduction half-reaction: O 2 (g) +2H 2 O(l) + 4e- → 4OH- (aq)  2Fe 2+ ( aq ) + 4OH- ( aq ) → 2Fe(OH) 2 (s)

15. Rusting Conditions Presence of water, air, and electrolytes (like salts) accelerate rusting on iron nails. Fe 2+ Fe 2+ Fe 2+ OH - OH - ½ O 2 H 2 O H 2 O H 2 O OH - OH - ½ O 2 H 2 O H 2 O H 2 O Fe 2+ Fe 2+ Fe 2+ H 2 O H 2 O H 2 O OH - OH - ½ O 2 H 2 O OH - OH - ½ O 2 H 2 O Cl - Na + Na + Cl - The salt ions enhance the ion exchange , speeding up both reactions. Cl⁻ ions, in particular, are aggressive and promote pitting corrosion by breaking down protective oxide layers.

16. Effect of Salt How It Accelerates Rusting ↑ Electrical conductivity Faster electron flow, speeds redox reactions Ion availability Promotes anodic and cathodic reactions Chloride ion action Breaks protective oxide layers, causes pitting corrosion Moisture attraction Keeps iron wet, prolongs exposure to corrosive conditions Prevents passivation Soluble rust doesn’t seal the surface, rusting continues Rusting Conditions

17. Electrochemistry of Rusting Anodic and cathodic sites form on the nail's surface, setting up redox reactions that produce rust. It is likely that the nail tip and the head had more defects than the center, so corrosion started at the tips.

18. Galvanic Corrosion Iron in contact with a more noble metal corrodes faster due to electron transfer. H 2 O H 2 O H 2 O H 2 O H 2 O H 2 O Fe 2+ Fe 2+ Fe 2+ Fe 2+ Fe 2+ Fe 2+ H 2 O H 2 O H 2 O H 2 O H 2 O H 2 O Cu 2+ Cu 2+ Cu 2+ Cu 2+ Cu 2+ Cu 2+ H 2 O H 2 O + + + + Larger circle = holds onto electrons more weakly Smaller circle = holds onto electrons more strongly Ions dissolve more readily Fewer ions dissolve More electrons on metal Fewer electrons on metal Potential for electrons to flow

19. Corrosion Prevention Fe 2+ Fe 2+ Fe 2+ OH - OH - ½ O 2 H 2 O OH - OH - ½ O 2 H 2 O Fe 2+ Fe 2+ Fe 2+ OH - OH - ½ O 2 H 2 O OH - OH - ½ O 2 H 2 O Mg 2+ Mg 2+ ½ H 2 OH - H 2 O ½ H 2 OH - H 2 O ½ H 2 OH - H 2 O ½ H 2 OH - H 2 O ½ H 2 OH - H 2 O ½ H 2 OH - H 2 O Mg 2+ Mg 2+ ½ H 2 OH - ½ H 2 OH - H 2 O OH - OH - ½ O 2 H 2 O H 2 O Mg 2+ ( aq ) + 2e - ⇌ Mg(s) -2.37 V H 2 O(l) + e - ⇌ ½H 2 (g) + OH - ( aq ) -0.83 V Fe 2+ ( aq ) + 2e - ⇌ Fe(s) -0.45 V ½O 2 (g) + H 2 O(l) +2e - ⇌ 2OH - ( aq ) +0.40 V

20. Galvanisation is the process of coating iron or steel with a layer of zinc (Zn) to protect it from rusting (corrosion). It is one of the most effective and widely used methods for corrosion prevention. Corrosion Prevention

21. Corrosion Prevention Barrier Protection Zinc acts as a physical shield that prevents moisture, oxygen, and corrosive agents from reaching the underlying iron or steel. Sacrificial Protection (Cathodic Protection) Even if the zinc coating is scratched or damaged, zinc continues to protect iron . Zinc is more reactive (more anodic) than iron and corrodes first , preventing iron from rusting: Zn → Zn 2+ + 2e− These electrons flow to the iron, inhibiting its oxidation .

22. Future of Nails Advancements include smart coatings, corrosion sensors, and eco-friendly manufacturing processes.

23. Biography The iron nail represents the meeting point of history, chemistry, culture, and engineering. Fe₂O ₃ + 3C → 2Fe + 3CO ↑ 4Fe+3O 2 ​+6H 2 ​O → 4Fe(OH) 3 ​→ Fe 2 ​O 3 ​⋅xH 2 ​O ( Death ). ( Life )

24. Thank You Questions and Discussion